Organic Chemistry/Foundational concepts of organic chemistry/Electron dot structures

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Electron Dot Structures

Electron dot structures, also called Lewis structures, give a representation of the valence electrons surrounding an atom.

Each valence electron is represented by one dot, thus, a lone atom of hydrogen would be drawn as an H with one dot, whereas a lone atom of Helium would be drawn as an He with two dots, and so forth.

Representing two atoms joined by a covalent bond is done by drawing the atomic symbols near to each other, and drawing a single line to represent a shared pair of electrons. It is important to note: a single valence electron is represented by a dot, whereas a pair of electrons is represented by a line.

The covalent compound hydrogen fluoride, for example, would be represented by the symbol H joined to the symbol F by a single line, with three pairs (six more dots) surrounding the symbol F. The line represents the two electrons shared by both hydrogen and fluorine, whereas the six paired dots represent fluorine's remaining six valence electrons.

Dot structures are useful in illustrating simple covalent molecules, but the limitations of dot structures become obvious when diagramming even relatively simple organic molecules. The dot structures have no ability to represent the actual physical orientation of molecules, and they become overly cumbersome when more than three or four atoms are represented.

Lewis dot structures are useful for introducing the idea of covalence and bonding in small molecules, but other model types have much more capability to communicate chemistry concepts.

Drawing electron dot structures

Some examples of electron dot structures for a few commonly encountered molecules from inorganic chemistry.

A note about Gilbert N. Lewis

Lewis was born in Weymouth, Massachusetts as the son of a Dartmouth-graduated lawyer/broker. He attended the University of Nebraska at age 14, then three years later transferred to Harvard. After showing an initial interest in Economics, Gilbert Newton Lewis earned first a B.A. in Chemistry, and then a Ph.D. in Chemistry in 1899.

For a few years after obtaining his doctorate, Lewis worked and studied both in the United States and abroad (including Germany and the Philippines) and he was even a professor at M.I.T. from 1907 until 1911. He then went on to U.C. Berkeley in order to be Dean of the College of Chemistry in 1912.

In 1916 Dr. Lewis formulated the idea that a covalent bond consisted of a shared pair of electrons. His ideas on chemical bonding were expanded upon by Irving Langmuir and became the inspiration for the studies on the nature of the chemical bond by Linus Pauling.

In 1923, he formulated the electron-pair theory of acid-base reactions. In the so-called Lewis theory of acids and bases, a "Lewis acid" is an electron-pair acceptor and a "Lewis base" is an electron-pair donor.

In 1926, he coined the term "photon" for the smallest unit of radiant energy.

Lewis was also the first to produce a pure sample of deuterium oxide (heavy water) in 1933. By accelerating deuterons (deuterium nuclei) in Ernest O. Lawrence's cyclotron, he was able to study many of the properties of atomic nuclei.

During his career he published on many other subjects, and he died at age 70 of a heart attack while working in his laboratory in Berkeley. He had one daughter and two sons; both of his sons became chemistry professors themselves.

Formal Charge

The formal charge of an atom is the charge that it would have if every bond were 100% covalent (non-polar). Formal charges are computed by using a set of rules and are useful for accounting for the electrons when writing a reaction mechanism, but they don't have any intrinsic physical meaning. They may also be used for qualitative comparisons between different resonance structures (see below) of the same molecule, and often have the same sign as the partial charge of the atom, but there are exceptions.

The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that "belong" to it in the Lewis structure when one counts lone pair electrons as belonging fully to the atom, while electrons in covalent bonds are split equally between the atoms involved in the bond. The total of the formal charges on an ion should be equal to the charge on the ion, and the total of the formal charges on a neutral molecule should be equal to zero.

For example, in the hydronium ion, H₃O⁺, the oxygen atom has 5 electrons for the purpose of computing the formal charge—2 from one lone pair, and 3 from the covalent bonds with the hydrogen atoms. The other 3 electrons in the covalent bonds are counted as belonging to the hydrogen atoms (one each). A neutral oxygen atom has 6 valence electrons (due to its position in group 16 of the periodic table); therefore the formal charge on the oxygen atom is 6 – 5 = +1. A neutral hydrogen atom has one electron. Since each of the hydrogen atoms in the hydronium atom has one electron from a covalent bond, the formal charge on the hydrogen atoms is zero. The sum of the formal charges is +1, which matches the total charge of the ion.

Formal Charge: number of valence electrons for an atom - (number of lone pair electrons + number electrons in bonds/2)

$FC=N_{valence}-\left(N_{lonepairs}+{\frac {N_{electrons}}{2}}\right)$

In chemistry, a formal charge (FC) on an atom in a molecule is defined as:

FC = number of valence electrons of the atom - ( number of lone pair electrons on this atom + total number of electrons participating in covalent bonds with this atom / 2).

$FC=N_{valence}-\left(N_{lonepairs}+{\frac {N_{participatingelectrons}}{2}}\right)$

When determining the correct Lewis structure (or predominant resonance structure) for a molecule, the structure is chosen such that the formal charge on each of the atoms is minimized.

Examples

carbon in methane

$FC=4-\left(0+{\frac {8}{2}}\right)=0$

Nitrogen in $NO_{2}^{-}$

$FC=5-\left(2+{\frac {6}{2}}\right)=0$

double bonded oxygen in $NO_{2}^{-}$

$FC=6-\left(4+{\frac {4}{2}}\right)=0$

single bonded oxygen in $NO_{2}^{-}$

$FC=6-\left(6+{\frac {2}{2}}\right)=-1$

Methane ($CH_{4}$): black is carbon, white is hydrogen

Nitrogen dioxide ($NO_{2}$): blue is nitrogen, red is oxygen

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